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# Chemistry

Chemistry. Chapter 3 Scientific Measurement. Qualitative Measurement. Gives results in a descriptive form Nonnumeric. Quantitative Measurement. Gives results in a definite form Usually as numbers and units. Scientific Notation. Shorthand way to express very large and very small numbers.

## Chemistry

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1. Chemistry Chapter 3 Scientific Measurement

2. Qualitative Measurement • Gives results in a descriptive form • Nonnumeric

3. Quantitative Measurement • Gives results in a definite form • Usually as numbers and units

4. Scientific Notation • Shorthand way to express very large and very small numbers

5. Example 3.6 x 104 = 3.6 x 10 x 10 x 10 x 10 = 36 000

6. 0.0081 = 8.1 x 10-3

7. Direction of decimal movement • To the left is +

8. To the right is -

9. Operations with numbers in scientific notation

10. Multiplication • Multiply the numbers and then add the exponents

11. Division • Divide the numbers and subtract the exponents

12. Addition and subtraction • Exponent must be the same to proceed

13. Must move the decimal appropriately and then adjust the exponent Then you can solve the problem

14. Measured values only as reliable as the instrument used to take the measurement!

15. Uncertainty in measurement • Accuracy • Measure of how close a measurement comes to the actual or true value

16. Precision • Measure of how close a series of measurements are to one another

17. Pg. 64 Dartboard example • In class: Pg. 97 #80

18. Evaluating the accuracy of a measurement • Percent error • Percent error = [error] X 100 accepted value

19. Error - the difference between the accepted value and the experimental value (absolute value)

20. Experimental value – measured in the lab • Accepted value – correct value based on reliable references • Pg. 65 example

21. Everyone understand so far? • Good!!!

22. Significant figures in measurements (sig figs) • Rules page 66-67

23. Sample problems Pg. 68

24. Sig Figs in Calculations • Rules for rounding Pg. 68 • Page 69 Sample

25. Solving problems with sig figs • Multiplying and dividing with sig figs • The answer you get must be rounded to the same number of sig figs as the measurement with the lowest number of sig figs (that you multiplied or divided)

26. Example • Multiply 4.610 feet by 1.7 feet. Express your answer in correct sig figs • 4.610 x 1.7 = 7.837 • How do you round it? • 4.610 has 4 sig figs • 1.7 has 2 sig figs • Round answer to 2 sig figs • Answer = 7.8 square feet

27. Adding and Subtracting with sig figs • When adding or subtracting measurements, the answer cannot have more certainly than the least certain measurement. • Answer must have the same number of sig figs to the right of the decimal point as the measurement with the fewest sig figs to the right of the decimal point

28. Example • 4.271 grams (3 sig figs to the right of decimal) • 2 grams (0 sig figs to the right of decimal) • + 10.0 grams (1 sig fig to the right of decimal) • 16.271 grams  round 16 grams

29. Handout practice – work with a partner! • Grab a calculator

30. SI System of Units • Page 73 Units of measurement • Table 3.1

31. Metric system established in France in 1790 • SI Adopted by international agreement in 1960

32. Prefixes • Page 74 Table 3.2

33. Length • SI unit - meter (m) • Pg. 74 Table 3.3

34. Volume • Space occupied by any sample of matter • L x W x H • “Derived” unit • Pg. 75 Table 3.4

35. Volume of a cube 1m on each side • SI unit = m3 • More common to use Liter (L) = dm3

36. 1 Liter • the volume occupied by a cube 10 cm on each side

37. 10 cm x 10 cm x 10 cm = 1000 cm3 1000 cm3 = 1 L

38. 1 dm = 10 cm

39. 1 L = 1 dm3

40. 1 mL = 0.001 L • 1000 mL = 1 L • 1000 cm3 = 1000 mL = 1 L

41. Much more dramatic with gases • Measuring devices calibrated at 20oC • Room temperature

42. Mass • Difference between mass and weight • SI unit = Kilogram (kg) • 1 g = 0.001 kg • 1000 g = 1 kg • Pg. 76 Table 3.5

43. Will show on board something special about H2O

44. Temperature Scales • Celsius • Kelvin Absolute zero

45. Kelvin scale explanation

46. Heat measurement • calorie • Joule • 1 cal = 4.184 J • 1J = 0.2390 cal

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